South Pasadena • AP Chemistry Name_____________________________________

Period _____ Date ___/___/___ 1•Matter and Measurement

1. How many

significant digits are

present in the

temperature read

from the

thermometer

illustrated to the

right?

a) 1 b) 2 c) 3 d) 4

2. The dimensions of a rectangular solid are 8.00 cm long, 4.00 cm wide, and 2.00 cm high. If the

density of the solid is 10.0 g/cm3, what is its mass?

a) 10/64 grams d) 320 grams

b) 10.0 grams e) 640 grams

c) 64.0 grams

3. A metal sample weighing 30.9232 grams was added to a graduated cylinder containing 23.26 mL

of water. The volume of water plus the sample was 24.85 mL. Which setup will result in the

density of this metal?

a) 30.9232 x (24.85-23.26)

b)

c)

d) 30.9232 x

e)

4. The number of significant digits in 0.30500 is

a) 1 d) 4

b) 2 e) 5

c) 3

5. A box measures 3.50 cm x 2.915 cm. The product of these numbers = 10.2025 cm2. What is the

proper way to report the area of the box?

a) 10.20 cm2 c) 10 cm2

b) 10.2 cm2 d) 10. cm2

6. The result of 2.350 x (4.0 + 6.311) is,

a) 24 c) 24.21

b) 24.2 d) 24.205

7. A student does a calculation using her calculator and the number 280.27163 is shown on the

display. If there are actually three significant figures, how should she show the final answer?

a) 280 d) 2.80 x 10-2

b) 280.3 e) 2.80 x 102

c) 280.27

8. The term that refers to the reproducibility of a laboratory measurement is

a) precision c) accuracy

b) repeatability d) exactness

9. Which measurement below is NOT written with three significant digits?

a) 2.00 cm c) 0.003 L

b) 550. grams d) 12.7 mm

10. The number 6.33 x 102 equals,

a) 6.33 c) 633

b) 0.633 d) 0.0633

11. All the following are characteristic properties of phosphorus. Which one is a chemical property?

a) Both red phosphorus and white phosphorus exist in solid allotropic forms.

b) The red form melts at about 600°C and the white form melts at 44°C.

c) The white form is soluble in liquid carbon disulfide, but is insoluble in water.

d) When exposed to air, white phosphorus will burn spontaneously, but red phosphorus will not.

12. Classify each observation as a physical or a chemical property and tally them.

Observation 1: Bubbles form on a piece of metal when it is dropped into acid.

Observation 2: The color of a crystalline substance is yellow.

Observation 3: A shiny metal melts at 650°C.

Observation 4: The density of a solution is 1.84 g/cm3

a) 2 chemical properties and 2 physical properties

b) 3 chemical properties and 1 physical properties.

c) 1 chemical properties and 3 physical properties

d) 4 chemical properties

e) 4 physical properties

13. Chromatography is a good way to separate the

a) elements in a compound

b) the components in a mixture

c) the atoms in an element

d) the phases of a pure substance

14. When a pure solid substance was heated, a student obtained another solid and a gas, each of which

was a pure substance. From this information which of the following statements is ALWAYS a

correct conclusion?

a) The original solid is not an element.

b) Both products are elements.

c) The original solid is a compound and the gas is an element.

d) The original solid is an element and the gas is a compound.

e) Both products are compounds.

15. The prefix “milli-” corresponds to what multiplication factor?

a) 10-6 d) 103

b) 10-3 e) 106

c) 101

16. A solution of sugar water may be defined as a

a) heterogeneous mixture

b) homogeneous mixture

c) heterogeneous compound

d) homogeneous compound

e) homogeneous element

17. “Wafting” is the proper technique for

a) neutralizing a spilled acid.

b) putting out burning clothing.

c) washing chemicals from the eye.

d) smelling a chemical substance.

e) observing the color of a chemical.

18. You measure the density of a slab of lead as 11.10 g/mL. The accepted value is 11.34 g/mL. The

percent error for your measurement is

a) 2.1 % c) 3.7 %

b) 2.4 % d) 5.1 %

19. Which one of the following elements is correctly matched with its symbol?

a) Ag, gold

b) Ni, nickel

c) Fl, fluorine

d) Mg, manganese

e) H, helium

Atoms and Elements

1. Certain properties are characteristic of metals. Which property means that you can pound the

substance into a foil?

a) ductility c) sectility

b) conductivity d) malleability

2. Which of the following is a metalloid?

a) As b) Ag c) S d) Pb e) He

3. Which of the following is a transition metal?

a) Cl b) Ni c) P d) Ca e) C

4. Which of the following is an alkali metal?

a) Mg b) Kr c) K d) Al e) H

5. Which of the following is an lanthanide?

a) Xe b) Eu c) Cd d) P e) W

6. Which element has the highest melting point?

a) Pb b) Au c) Os d) W e) Hg

7. Cathode rays start at the

a) negative electrode c) positive electrode

b) power source d) gas inside the tube

8. In a cathode ray tube, electrons are bent toward

a) a positively charged plate.

b) a negatively charged plate.

9. Listed below are the charges and masses of four particles. Which one will be deflected the least in

a mass spectrometer?

a) +2, 2 amu c) +1, 1 amu

b) +4, 4 amu d) +1, 4 amu

10. In a Millikan oil drop type experiment, the charge on four oil drops (in Coulombs) was found to be:

3.33 Coulombs

8.88 Coulombs

6.66 Coulombs

11.10 Coulombs

What is the charge on the electron according to this experiment?

a) 1.11 Coulomb c) 4.44 Coulomb

b) 2.22 Coulomb d) 11.10 Coulomb

11. Pictured below is a schematic of the Rutherford experiment. Which scattered -particle gives the best

evidence for the nuclear atom?

a) a b) b c) c d) d e) e

12. Which of the following is an isotope of the element with 20 protons (p=20) and 22 neutrons (n=22)?

a) titanium-22 c) calcium-40

b) zirconium-40 d) titanium-48

13. The imaginary element X has the following natural abundances and isotopic masses. What is the

atomic mass of X?

X 24.02 amu 40.0%

X 26.10 amu 60.0%

Show your work:

For questions 14 - 17, use the following key:

(each answer may be used once, more than once,

or not at all)

a) alpha

b) beta

c) gamma

d) alpha and beta, but not gamma

14. A high energy form of light

15. Two protons & two neutrons

16. A high speed electron

17. Used by Ernest Rutherford as a “probe”

For questions 18 - 22, use the following key:

(each answer may be used once, more than once,

or not at all.)

a) John Dalton

b) Ernest Rutherford

c) J.J. Thomson

d) Democritus

18. His model of the atom has been called the “plum pudding” Model.

19. His model of the atom has been called the “billiard ball” model.

20. He studied matter in cathode ray tubes.

21. His philosophical idea included the term “atomos”.

22. He added to the atomic theory the idea that atoms had positive and negative parts.

23. Consider the following notation: Rn

Which statement below is correct?

a) This particle contains 86 protons

b) This particle has a mass number of 86

c) This particle has an atomic number of 220

d) This particle contains 220 neutrons

25. If copper metal is a mixture two isotopes, Cu-63, mass = 62.9298 u and Cu-65, mass = 64.9278 u. The

molar mass of copper is 64.546 g/mole. Calculate the % abundances of the two isotopes of copper.

Show your work.

Just For Fun:

Element names finish these sentences.

• A ridiculous inmate is a ___.

• I bumped my ___ the car door.

• I am sad when all the flowers ____.

• What the police officer does to the crook. ___

• What the doctor does to the patient. ___

• What the undertaker does if the doctor doesn’t succeed. ___

• If your cattle get away, ___.

• A famous London theatre is the ___.

• Demonstrations help keep the lectures from getting ___.

Molecules and Compounds1. What is the formula of the ionic compound formed between Mg and Br?

a) MgBr d) Mg2Br2

b) Mg2Br e) Mg2Br3

c) MgBr2

2. What is the formula of the ionic compound formed between Ca and P?a) Ca2P3 d) Ca2P

b) CaP e) Ca3P2

c) Ca5P10

3. What is the name of the SO32– ion?

a) sulfate d) sulfur trioxide

b) nitrate e) hydrogen sulfate

c) sulfite

4. What is the correct formula and charge for the chromate ion?a) CrO4

2– d) Cr2O7–

b) CrO4– e) Cr3+

c) Cr2O72–

5. Which one of the following elements forms ions with two different valences?

a) calcium c) iron

b) arsenic d) fluorine

6. The correct name for CCl4 is

a) carbon(I) chloride

b) carbon chloride

c) carbon tetrachloride

d) monocarbon chloride(IV)

e) carbochlorinate

7. The correct formula for hydrogen telluride is

a) HTe c) H3Te

b) H2Te d) HTe2

8. The correct formula for dinitrogen tetroxide isa) NO2 d) NO3

–

b) N2O4 e) (N2O)4

c) N2O5

9. The correct name for S2Cl2 is

a) sulfur dichloride

b) sulfur(I) chloride

c) sulfur(II) chloride

d) disulfur dichloride

e) sulfur chloride

10. The correct name for NO2 is

a) nitrogen dioxide

b) nitrite

c) nitrogen oxide

d) nitrogen(II) oxide

e) nitrate

11. The molar mass of (NH4)2S is closest to:

a) 50 g/mol c) 68 g/mol

b) 82 g/mol d) 100 g/mol

12. How many atoms are in 12 molecules of glucose, C6H12O6?

a) 24 c) 2160

b) 288 d) 7.22 x 1024

13. Calculate the number of atoms in 4.0 x 10-5 g of aluminum.

a) 8.9 x 1017 c) 6.5 x 1020

b) 4.6 x 1019 d) 3.8 x 1023

14. Which of the following samples contains the smallest number of atoms?

a) 1 g H2 c) 1 g O3

b) 1 g O2 d) 1 g Cl2

15. What is the mass of one molecule of octane, C8H18?

a) 114 g c) 1.10 x 10-22 g

b) 1.89 x 10-22 g d) 4.32 x 10-23 g

16. What is the percent nitrogen (by mass) in ammonium carbonate, (NH4)2CO3?

a) 14.53% c) 29.16%

b) 27.83% d) 33.34%

17. Of the following, the only empirical formula isa) N2F2 c) H2C2

b) N2F4 d) HNF2

18. A compound consists of the following elements by weight percent:

carbon - 40.0%

oxygen - 53.3%

hydrogen - 6.7%

The ratio of carbon : oxygen : hydrogen in the empirical formula is

a) 1:2:1 c) 1:1:2

b) 1:1:1 d) 2:1:2

19. An organic compound which has the empirical formula CHO has a molar mass of 232. Its molecular

formula is:

a) CHO c) C4H4O4

b) C2H2O2 d) C8H8O8

20. When CaSO4·y H2O is heated, all of the water is driven off. If 34.0 g of CaSO4 [molar mass = 136] is

formed from 43.0 g of CaSO4·y H2O, what is the value of y?

a) 1 c) 3

b) 2 d) 4

Chemical Equations and Stoichiometry

1. Balance the following equation:___NH3 + ___O2 ___NO2 + ___H2O

The balanced equation shows that 1.00 mole of NH3 requires ___ mole(s) of O2.

a) 0.57 c) 1.33b) 1.25 d) 1.75

2. Write a balanced equation for the combustion of acetaldehyde, CH3CHO.

When properly balanced, the equation indicates that ___ mole(s) of O2 are required for each mole of

CH3CHO.

a) 1 c) 2.5b) 2 d) 3

3. Balance the following equation with the SMALLEST WHOLE NUMBER COEFFICIENTS possible. Select the number that is the sum of the coefficients in the balanced equation:

___KClO3 ___KCl + ___O2a) 5 b) 6 c) 7 d) 8

4. Write a balanced equation for the combustion of propane, C3H8.

When properly balanced, the equation indicates that ___ moles of O2 are required for each mole of

C3H8.

a) 3 b) 3.5 c) 5 d) 8

5. What is the total mass of products formed when 16 grams of CH4 is burned with excess oxygen?a) 80 g c) 36 gb) 44 g d) 32 g

6. Calculate the mass of hydrogen formed when 25 g of aluminum reacts with excess hydrochloric acid.2Al + 6HCl 2 AlCl3 + 3 H2

a) 0.41 g c) 1.2 gb) 0.92 g d) 2.8 g

7. How many grams of the mixed oxide, Fe3O4, are formed when 6.00 g of O2 react with Fe according

to3Fe + 2O2 Fe3O4

a) 43.4 c) 174b) 86.8 d) 21.7

8. For the reaction:2MnO2 + 4KOH + O2 + Cl2 2KMnO4 + 2KCl + 2H2O

there is 100. g of each reactant available. Which reagent is the limiting reagent?[Molar Masses: MnO2=86.9; KOH=56.1; O2=32.0; Cl2=70.9]

a) MnO2 c) O2b) KOH d) Cl2

9. How many grams of nitric acid, HNO3, can be prepared from the reaction of 92.0 g of NO2 with

36.0 g H2O?

3NO2 + H2O 2HNO3 + NO

a) 64 c) 84b) 76 d) 116

10. The reaction of 25.0 g benzene, C6H6, with excess HNO3 resulted in 21.4 g C6H5NO2. What is the

percentage yield?C6H6 + HNO3 C6H5NO2 + H2O

a) 100% c) 54.3%b) 27.4% d) 85.6%

11. How many grams of H2O will be formed when 16.0 g H2 is allowed to react with 16.0 g O2

according to2H2 + O2 2H2O?

a) 18.0 g c) 9.00 gb) 144 g d) 32.0 g

12. When 8.00 g of H2 reacts with 32.0 g of O2 in an explosion, 2H2 + O2 2H2O, the final gas mixture

will contain:a) H2, H2O, and O2 c) O2 and H2O only

b) H2 and H2O only d) H2 and O2 only

13. 1.056 g of metal carbonate, containing an unknown metal, M, were heated to give the metal oxide and 0.376 g CO2.

MCO3(s) + heat MO(s) + CO2(g)

What is the identity of the metal M?a) Mg c) Znb) Cu d) Ba

14. Styrene, the building block of polystyrene, is a hydrocarbon, a compound containing only C and H. A given sample is burned completely and it produces 1.481 g of CO2 and 0.303 g of H2O. Determine

the empirical formula of the compound.a) CH c) C2H3b) CH2 d) C2H5

Reactions In Aqueous Solution

1. On the basis of the solubility rules, which of the following is insoluble?a) K2O d) (NH4)2SO4

b) Na2CO3 e) Ba(C2H3O2)2

c) PbS

2. In a double replacement reaction, formation of which of the following does not necessarily lead to

a chemical change?a) HC2H3O2 d) H2S

b) AgCl e) NaClc) CO2

3. Reaction of an acid with a carbonate (such as CaCO3) always results in the formation of

a) O2 d) O3

b) C(diamond) e) CO2

c) CH4

4. Which of the following is incorrect?a) all salts containing NH4

+ are soluble.

b) all salts containing NO3– are soluble.

c) all fluorides are soluble.

d) all sulfates (except those of Ca2+, Sr2+, Ba2+, and Pb2+) are soluble.e) most hydroxides are insoluble, except those of Ca2+, Sr2+, Ba2+, the alkali metals and NH4+.

5. One of the gases shown below is NOT usually formed in a double replacement reaction. Which

one?a) N2 d) NH3

b) CO2 e) H2S

c) SO2

6. Write the balanced molecular equation for the reaction of washing soda, Na2CO3 and vinegar,

HC2H3O2.

7. The net ionic equation for the above reaction is:

8. How many moles of H+ are associated with the acid, H2SO3, during neutralization?

a) 0 b) 1 c) 2 d) 3

9. How many moles Al2O3 are needed to neutralize 1 mole of HCl?

a) 1/3 d) 6

b) 2/3 e) 12

c) 2 f) 1/6

10. Write the net reaction that will occur when solid ammonium carbonate is added to a solution of

hydrosulfuric acid.

11. When H2SO4 and Ba(OH)2 are reacted in a double replacement reaction, one of the products of the

reaction is…a) H2 d) BaH2

b) H2O e) SO2

c) BaS

12. In the double replacement reaction between the weak acid, HC2H3O2 and strong base, NaOH,

which ion(s) are spectator ions?a) Na+, C2H3O2

– d) H+, C2H3O2–

b) Na+, OH– e) Na+ only

c) OH– only

13. Which of the following is a base?a) KOH d) CH3OH

b) C2H5OH e) CO2

c) Br–

14. Which of the following is a strong acid?a) H2CO3 d) HClO3

b) HF e) HNO3

c) H3PO4

15. Which of the following is an acid in aqueous solutions?a) H2CO3 d) H2O

b) Al2O3 e) BaO

c) CH4

16. SO2 turns into which acid in solution?

a) HNO3 d) H2S

b) H2SO3 e) HNO2

c) H2SO4

17. What is the oxidation number of C in CO32–?

a) +6 d) +1

b) +4 e) –1

c) +2

18. What is the oxidation number of Br in KBrO4?

a) +1 b) –1 c) +5 d) +7 e) +8

19. For each change below, label the change of the underlined element as Oxidation, Reduction, or

Neither

___ Cu2+ Cu___ CH4 CO2

___ H2O2 H2O

___ CO2 H2CO3

20. How many milliliters of 0.123 M NaOH solution contain 25.0 g of NaOH (molar mass = 40.00

g/mol)?

a) 5.08 mL d) 625 mL

b) 50.8 mL e) 5080 mL

c) 508 mL

21. If you need 1.00 L of 0.125 M H2SO4, how would you prepare this solution?

a) Add 950. mL of water to 50.0 mL of 3.00 M H2SO4.

b) Add 500. mL of water to 500. mL of 0.500 M H2SO4.

c) Add 750 mL of water to 250 mL of 0.375 M H2SO4.

d) Dilute 36.0 mL of 1.25 M H2SO4 to a volume of 1.00 L.

e) Dilute 20.8 mL of 6.00 M H2SO4 to a volume of 1.00 L.

22. What is the ion concentration in a 0.12 M solution of BaCl2?

a) [Ba2+] = 0.12 M and [Cl] = 0.12 M.

b) [Ba2+] = 0.12 M and [Cl] = 0.060 M.

c) [Ba2+] = 0.12 M and [Cl] = 0.24 M.

d) [Ba2+] = 0.060 M and [Cl] = 0.060 M.e) [Ba+] = 0.12 M and [Cl2] = 0.12 M.

23. What is the molarity of the solution that results when 60.0 g NaOH is added to enough water to

make 500. mL solution?

a) 1.33 M d) 8.0 M

b) 12.0 M e) 1.50 M

c) 3.00 M

24. What is the molarity of the solution that results when 45.0 g HCl is dissolved in enough water to

make 250. mL solution?

a) 4.94 M d) 1.80 M

b) 4.50 M e) 1.46 M

c) 3.24 M

25. What is the concentration of Cl– ion in 0.60 M AlCl3 solution?

a) 1.8 M d) 0.30 M

b) 0.60 M e) 0.10 M

c) 0.20 M

26. How many grams of Na2CO3 (molar mass = 106.0 g/mol) are required for complete reaction with

25.0 mL of 0.155 M HNO3?

Na2CO3 + 2HNO3 2NaNO3 + CO2 + H2O

a) 0.122 g d) 20.5 g

b) 0.205 g e) 205 g

c) 0.410 g

27. What volume of 0.150 M NaOH is needed to react completely with 3.45 g iodine according to the

equation:3 I2 + 6 NaOH 5 NaI + NaIO3 + 3 H2O

a) 181 mL d) 2.04 mL

b) 45.3 mL e) 1.02 mL

c) 4.08 mL

28. What is the concentration of an NaOH solution if it takes 16.25 mL of a 0.100 M HCl solution to titrate

25.00 mL of the NaOH solution?

a) 0.0165 M d) 0.100 M

b) 0.151 M e) 0.413 M

c) 0.0650 M

29. A 4.00 M solution of H3PO4 will contain ___g of H3PO4 in 0.250 L of solution.

a) 196 g d) 24.0 g

b) 98.0 g e) 12.0 g

c) 49.0 g

Energy and Chemical Reactions

1. How many joules are equivalent to 37.7 cal?

a) 9.01 J c) 1.51 J

b) 4.184 J d) 158 J

2. The quantity of heat that is needed to raise the temperature of a sample of a substance 1.00 degree is

called its

a) heat capacity c) enthalpy

b) specific heat d) kinetic energy

3. Equal masses of two substances, A & B, each absorb 25 Joules of energy. If the temperature of A

increases by 4 degrees and the temperature of B increases by 8 degrees, one can say that

a) the specific heat of A is double that of B.

b) the specific heat of B is double that of A.

c) the specific heat of B is negative.

d) the specific heat of B is triple that of A.

4. If 25 J are required to change the temperature of 5.0 g of substance A by 2.0C, what is the specific

heat of substance A?

a) 250 J/gC c) 10. J/gCb) 63 J/gC d) 2.5 J/gC

5. How much energy is required to change the temperature of 2.00 g aluminum from 20.0C to 25.0C?

The specific heat of aluminum is 0.902 J/gC.

a) 2.3 J c) 0.36 J

b) 9.0 J d) 0.090 J

6. Consider the thermal energy transfer during a chemical process. When heat is transferred to the

system, the process is said to be _______ and the sign of H is ________.

a) exothermic, positive

b) endothermic, negative

c) exothermic, negative

d) endothermic, positive

7. What is the E for a system which has the following two steps:

Step 1: The system absorbs 60 J of heat

while 40 J of work are performed on it.

Step 2: The system releases 30 J of heat

while doing 70 J of work.

a) 100 J c) 30 J

b) 90 J d) zero

8. When two solutions react the container “feels hot.” Thus,

a) the reaction is endothermic.

b) the reaction is exothermic.

c) the energy of the universe is increased.

d) the energy of both the system and the surroundings is decreased.

9. The equation for the standard enthalpy of formation of N2O3 is

a) N2O(g) + O2(g) N2O3(g)

b) N2O5(g) N2O3(g) + O2(g)

c) NO(g) + NO2(g) N2O3(g)

d) N2(g) + 3/2 O2(g) N2O3(g)

10. For the general reaction

2 A + B2 2 AB, H is +50.0 kJ.

We can conclude that

a) the reaction is endothermic.

b) the surroundings absorb energy.

c) the standard enthalpy of formation of AB is -50.0 kJ.

d) the molecule AB contains less energy than A or B2.

11. Calculate the enthalpy of combustion of C3H6:

C3H6(g) + 9/2O2(g) 3CO2 + 3H2O

using the following data:

3C(s) + 3H2(g) C3H6(g) H= 53.3 kJ

C(s) + O2(g) CO2(g) H=-394 kJ

H2(g) + 1/2O2(g) H2O(l) H=-286 kJ

a) -1517 kJ c) -626 kJ

b) 1304 kJ d) -2093 kJ

12. Which one of the following would have an enthalpy of formation value (Hf) of zero?

a) H2O(g) c) H2O(l)

b) O(g) d) O2(g)

13. Calculate the heat of vaporization of titanium (IV) chloride: TiCl4(l) TiCl4(g)

using the following enthalpies of reaction:

Ti(s) + 2Cl2(g) TiCl4(l) H=-804.2 kJ

TiCl4(g) 2Cl2(g) + Ti(s) H= 763.2 kJ

a) -1567 kJ c) 1165 kJ

b) -783.7 kJ d) 41 kJ

14. Calculate the enthalpy of reaction for:

D + F G + M

using the following equations and data:

G + C A + B H = +277 kJ

C + F A H = +303 kJ

D B + M H = -158 kJ

a) -132 kJ c) +422 kJ

b) -422 kJ d) +132 kJ

15. Calculate the standard enthalpy of the reaction for the process

3NO(g) N2O(g) + NO2(g)

using the standard enthalpies of formation (in kJ/mol): NO = 90; N2O = 82.1; NO2 = 34.0

a) -153.9 kJ c) -26.1 kJ

b) 206 kJ d) 386 kJ

16. The standard molar enthalpy of combustion is

-1277.3 kJ for the combustion of ethanol.

C2H5OH(l) + 3O2(g) 2CO2(g) + 3H2O(g)

Calculate the standard molar enthalpy of formation for ethanol based on the following standard

enthalpies of formation:

Hf CO2 = -393.5 kJ/mol

Hf H2O = -241.8 kJ/mol

a) -642.7 kJ/mol c) 235.1 kJ/mol

b) -235.1 kJ/mol d) 642.7 kJ/mol

17. Calculate the amount of heat needed to change 25.0 g ice at 0C to water at 0C.

The heat of fusion of H2O = 333 J/g;

a) 56.5 kJ c) 7.06 kJ

b) 8.33 kJ d) 463 kJ

7&8 • Atomic Structure & Periodicity

A = 2.18 x 10-18 J h = 6.626 x 10-34 J·s

R = 1.097 x 107 m–1 c = 3.00 x 108 m·s–1

mass of an electron = 9.11 x 10-31 kg

1. What wavelength corresponds to a frequency of 8.22 x 109 Hz?

a) 0.307 m d) 0.110 m

b) 0.0365 m e) 27.4 m

c) 0.122 m

2. A radio station transmits at 110 MHz (110 x 106 Hz). What wavelength is this radio wave?

a) 3.65 x 10–5 m c) 3.81 x 10–5 m

b) 3.30 m d) 2.73 m

3. Which one of the following is NOT a proper unit for frequency?

a) Hz c) m·s–1

b) s–1 d)

4. Calculate the wavelength of the fourth line in the Balmer series (the visible series) of the hydrogen

spectrum.

a) 0.12334 m d) 4.1029 x 10–7 m

b) 24.373 m e) 36.559 m

c) 2.7353 x 10–7 m

5. What is the relationship between the energy of a photon of light and its frequency?

a) E = d) E =

b) E = e) E =

c) E = h

6. What is the energy needed to raise an electron in the hydrogen atom from the second energy level to

the third energy level?

a) 1.52 x 104 J d) 4.48 x 10–19 J

b) 3.63 x 10–19 J e) 3.03 x 10–19 J

c) 2.18 x 10–19 J

7. What is the de Broglie wavelength of an electron moving at 80.0% the speed of light.

a) 3.03 x 10–12 m c) 3.30 x 1011 m

b) 2.42 x 10–12 m d) 1.59 x 10–25 m

9. Which quantum number determines the subshell occupied by an electron (s, p, d, f, etc.)?

a) n c) ml

b) l d) ms

10. What position on the standing wave shown below corresponds to a crest?

a) A b) B c) C d) D e) E

11. How many orbitals make up the 4d subshell?

a) 0 b) 1 c) 3 d) 5 e) 7

12. The value of l that is related to the following orbital is:

a) 0 b) 1 c) 2 d) 3 e) 4

13. The correct electron configuration for nitrogen is

a) 1s2 2s2 2p6 3s2 3p2

b) 1s2 2s2 2p6 2d4

c) 1s2 2s2 2p3

d) 1s2 2s2 3s2 4s1

e) 1s2 1p5

14. The electron configuration of the indicated atom in the ground state is correctly written for which

atom?

a) Ga [Ar] 3d12 4s2

b) Ni [Ar] 3d10

c) Ni [Ar] 3s2 3p8

d) Cu [Ar] 3d10 4s1

15. Which of the following sets of quantum numbers is possible for a 3d electron?

a) n = 3, l = 3, ml = –2, ms = +

b) n = 2, l = 1, ml = +1, ms = –

c) n = 3, l = 1, ml = 0, ms = –

d) n = 3, l = 2, ml = –2, ms = +

e) n = 4, l = 1, ml = +1, ms = +

16. In what section of the periodic table is the 4f subshell being filled?

a) period 4

b) transition elements Y to Cd

c) noble gases

d) group IA

e) lanthanides

17. Which one of the following elements has 3 electrons in a p subshell?

a) Sb b) Na c) Sc d) V e) Nd

18. Which of the following distributions of electrons is correct for three electrons in p-subshell?

a) b

)

c) d

)

e)

19. Which of the following particles would be most paramagnetic?

a) P

b) Ga

c) Br

d) Cl-

e) Na+

20. Which of the following correctly represents the ionization of an atom?

a) Cl(g) + e– Cl–(g)

b) Na(g) Na+(g) + e–

c) Na(s) – e– Na+(g)d) Cl2(g) 2 Cl(g)

21. Which of the following is likely to have the largest atomic radius?

a) H b) Mn c) Cl d) Rb e) Ag

22. Which one of the following isoelectronic species has the smallest radius?

a) Mg2+ d) F –

b) Na+ e) O2–

c) Ne

NOTE: explain your reasoning on the last page.

23. Which of the following has the greatest ionization energy?

a) K b) Ca c) Fe d) Ga e) Br

24. Which of the following has the lowest ionization energy?

a) Li b) Na c) K d) Rb e) Cs

25. The successive ionization energies for one of the period three elements is listed below. Which

element is referred to?

E1 577.4 kJ/mol

E2 1,816 kJ/mol

E3 2,744 kJ/mol

E4 11,580 kJ/mol

E5 15,030 kJ/mol

a) Na b) Mg c) Al d) Si e) P

NOTE: explain your reasoning on the last page.

Write the electron configuration for silver:

27. Long form:

28. Short form:

29. Explain your answer to question 22.

Which one of the following isoelectronic species has the smallest radius?

a) Mg2+

b) Na+

c) Ne

d) F –

e) O2–

30. Explain your answer to question 25.

The successive ionization energies for one of the period three elements is listed below. Which

element is referred to?

E1 577.4 kJ/mol

E2 1,816 kJ/mol

E3 2,744 kJ/mol

E4 11,580 kJ/mol

E5 15,030 kJ/mol

a) Na b) Mg c) Al d) Si e) P

Bonding & Molecular Structure1. The correct Lewis symbol for ground state carbon is

a) b) c) d) e)

2. The correct Lewis symbol for ground state aluminum is

a) b) c) d) e)

3. Using the picture below, what process corresponds to the lattice energy?

a) 1 b) 2 c) 3 d) 4 e) 5

4. Which of the following favors formation of an ionic compound?

a) low ionization energy for metal

b) high electron affinity for non metal

c) high lattice energy

d) all of a-c above

e) none of a-c above

5. Which of the atoms below is least likely to violate the octet rule?

a) Be b) P c) S d) B e) F

6. How many electrons are shown in the Lewis structure of perchlorate ion, ClO4–?

a) 30 b) 31 c) 32 d) 50 e) 51

Questions7 - 9 refer to the following energy diagram:

7. What is the energy of the two isolated atoms?

a) –400 J d) 0.100 nm

b) –335 J e) –155 J

c) 0 J

8. What is the bond length of the bond between the two atoms?

a) 0.020 nm d) 0.330 nm

b) 0.140 nm e) 2.00 nm

c) 0.400 nm

9. What is the bond energy (bond strength) of the bond?

a) 0 J d) –330 J

b) –110 J e) –422 J

c) –10 J

10. What is the bond order of the C–O bond in

a) 4 b) 1.5 c) 0.5 d) 1 e) 2

11. Which of the following is the correct Lewis structure for SOCl2? (Consider formal charge)

a) d)

b) e)

c)

12. As the bond order of a carbon-carbon bond increases, which one of the following decreases?

a) # of electrons between the carbon atoms

b) vibrational frequency of bond vibrations

c) bond energy (bond strength)

d) bond length

13. In which of the following is the actual compound a resonance hybrid of Lewis structures?

a) NO2 d) CCl4

b) H2O e) none of these

c) O3

14. Which of the following bonds is most polar?

a) N – Cl d) Br – Br

b) C – N e) S – O

c) S – S

15. Which one of the following molecules is a polar molecule?

a) c)

b) d)

16. Which of the following molecular shapes has six atoms joined to a central atom?

a) linear d) trigonal bipyramid

b) tetrahedral e) planar triangular

c) octahedral

17. Which molecular shape has bond angles which are not all the same?

a) linear d) planar triangular

b) tetrahedral e) trigonal bipyramid

c) octahedral

19. The molecule BrF3 has how many lone pairs of electrons on the central atom?

a) 0 b) 1 c) 2 d) 3

20. What is the geometrical arrangement of electron pairs in H2O?

a) linear d) trigonal bipyramidal

b) bent e) tetrahedral

c) octahedral

21. What is the shape of BrI3?

a) square planar d) pyramidal

b) T-shaped e) bent

c) distorted tetrahedral

22. What is the shape of the IF4– ion?

a) square planar d) octahedral

b) tetrahedral e) T-shaped

c) square pyramidal

23. Which of the following is a polar species?a) CO2

b) PCl5

c) ICl2–

d) TeCl4

e) CCl4

24. Among those listed below, which element will have the strongest tendency to form double bonds?

a) S b) B c) Al d) O

H2.1

Electronegativity Values He---

Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Ne---

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

Ar---

K0.8

Ca1.0

Ga1.6

Ge1.8

As2.0

Se2.4

Br2.8

Kr---

Rb0.8

Sr1.0

In1.7

Sn1.8

Sb1.9

Te2.1

I2.5

Xe---

Cs0.7

Ba0.9

Tl1.8

Pb1.8

Bi1.9

Po2.0

At2.2

Rn---

Fr0.7

Orbital Hybridization & Molecular Orbitals

1. Which hybridization is associated with a steric number of 3?

a) sp d) sp3d

b) sp2 e) sp3d2

c) sp3

2. The molecule BrF3 has a steric number of ___ on the central atom?

a) 3 b) 4 c) 5 d) 6

3. What is the hybridization of Br in BrF3?

a) sp d) sp3d

b) sp2 e) sp3d2

c) sp3

4. How many equivalent sp3d orbitals are there?

a) 3 b) 1 c) 5 d) 6

5. What type of hybridization is associated with a square planar molecular shape?

a) sp3 d) sp3d

b) sp2 e) sp3d2

c) sp

6. What shape for electron pairs is associated with sp3d2 hybridization?

a) linear d) tetrahedral

b) square planar e) octahedral

c) bent

7. What hybridization is predicted for phosphorus in the PCl3 molecule?

a) sp2 c) sp

b) sp3 d) sp3d2

8. A double bond contains ___ sigma bond(s)

and ___ pi bond(s).

a) 0, 2 b) 1, 2 c) 2, 0 d) 1, 1

9. What angle exists between orbitals in sp3d2 hybrid orbitals?

a) 90.0° d) 120.0°

b) 180.0° e) 78.5°

c) 109.5°

10. Which of the following elements is most likely to display sp3d hybridization?

a) oxygen d) carbon

b) nitrogen e) boron

c) phosphorus

11. How many sigma () and pi () electrons pairs are in a carbon dioxide molecule?

a) four and zero d) two and four b) three and two e) one and three c) two and two

12. What is the hybridization of the oxygen atoms in CH3OH and CO2, respectively?

a) sp3, sp3 d) sp2, sp2

b) sp3, sp2 e) sp3, sp

c) sp2, sp3

13. All of the following species contain two -bonds EXCEPT

a) SCN d) OCS

b) CO e) NO

c) H2CCO

14. Which response contains all the characteristics that should apply to BF3?

1. trigonal planar

2. one unshared pair of electrons on B

3. sp2 hybridized boron atom

4. polar molecule

5. polar bonds

a) 2, 4, and 5 d) 1, 3, and 5

b) 1, 3, and 4 e) 3, 4, and 5

c) 1, 2, and 3

Short Answer:15. Consider the structural formula for acetic acid, HC2H3O2 or CH3COOH. Indicate the type of

hybridization used by each of the carbon and oxygen atoms.

16. Consider the structural formula for the acetate ion, C2H3O2– or CH3COO–. Indicate the hybridization

used by each of the carbon and oxygen atoms.

17. What shape of bond is formed in the above ion? ______________

19. When BF3 + NH3 BF3NH3, how does the hybridization of the boron atom change, if at all?

Gases and Their Properties

1. A pressure of 745 mmHg corresponds to ___ kPa.

a) 55.89 kPa c) 99.3 kPa

b) 0.980 kPa d) 745 kPa

2. Liquid nitrogen has a boiling point of -196 °C this corresponds to…

a) - 469 K c) 153 K

b) 77 K d) 469 K

3. 1.20 atm is the same pressure as:

a) 1.2 mmHg d) 850 mmHg

b) 760 mmHg e) 358 mmHg

c) 912 mmHg

4. For an ideal gas, which pair of variables are inversely proportional to each other (if all other factors

remain constant)?

a) P, V c) V, T

b) P, T d) n, P

5. A real gas would act most ideal at

a) 1.0 atm and 273 K

b) 10 atm and 546 K

c) 10 atm and 273 K

d) 0.5 atm and 546 K

e) 0.5 atm and 273 K

6. One mole of hydrogen, H2, occupies 61.2 L at

a) 100 C and 1.00 atm

b) 200 C and 1.00 atm

c) 0 C and 0.500 atm

d) 50 C and 0.500 atm

e) 100 C and .500 atm

7. A 31.0 mL sample of gas is collected at a temperature of 37 °C and pressure of 720 mmHg. What

is its volume at 17 °C and 580 mmHg.

a) 23 mL d) 41 mL

b) 27 mL e) 58 mL

c) 36 mL

8. The coldest possible temperature of a gas is:

a) 0 °C b) 273 K c) -273 K d) -273 °C

9. The pressure of 4.0 L of an ideal gas in a flexible container is decreased to one-third of its original

pressure and its absolute temperature is decreased by one-half. The volume then is

a) 1.0 L b) 4.0 L c) 6.0 L d) 8.0 L e) 24 L

10. A given mass of gas in a rigid container is heated from 100 C to 300 C. Which of the following

best describes what will happen to the pressure of the gas? The pressure will…

a) decrease by a factor of three.

b) increase by a factor of three.

c) increase by a factor less than three.

d) decrease by a factor greater than three.

11. What is the pressure exerted by some nitrogen gas collected in a tube filled with water on a day

when the room temperature is 18.0 °C and the room pressure is 750.0 mmHg? [The partial pressure

of water at 18 °C is 15.5 mmHg.]

a) 15.5 mmHg d) 760.0 mmHg

b) 750.0 mmHg e) 732.0 mmHg

c) 734.5 mmHg

12. As the average kinetic energy of the molecules of a sample increases, the temperature of the sample

a) decreases c) remains the same

b) increases

13. If a gas that is confined in a rigid container is heated, the pressure of the gas will…

a) increase c) remain the same

b) decrease

14. A mixture of gases at 810 kPa pressure contains:

3.0 moles of oxygen gas,2.0 moles of helium gas, and4.0 moles of carbon dioxide gas.

What is the partial pressure of helium gas, PHe.

a) 405 kPa d) 81.0 kPa

b) 1620 kPa e) 180 kPa

c) 810 kPa

15. If a gas has a pressure of 2.0 atm, which one of the following equations will express its pressure

after...

• the number of moles has been increased to three times the original amount,

• the absolute temperature (K) has been reduced to half, and

• the volume has been tripled?a) P2 = 2.0 atm x x x

b) P2 = 2.0 atm x x x

c) P2 = 2.0 atm x x x

d) P2 = 2.0 atm x x x

16. A sample of gas occupies 30.0 L at 0.800 atm and 298 K. How many moles of gas are in the

sample?

a) 22.4 b) 0.981 c) 1.02 d) 2.23

e) none of these

17. When ammonium nitrite undergoes decomposition, only gases are produced according to the equation: NH4NO2(s) N2(g) + 2H2O(g)

What is the total volume of gases produced at 819K and 1.00 atm pressure when 128 g of

ammonium nitrite undergoes the above decomposition reaction?

18. At STP, it was found that 1.12 L of a gas had a mass of 2.78 g. Its molar mass is

a) 2.78 g/mol c) 55.6 g/mol

b) 27.8 g/mol d) 111 g/mol

19. A mixture of gases, nitrogen, oxygen, and carbon dioxide at 27 C and 0.50 atmospheres pressure

occupied a volume of 492 mL. How many moles of gas are there in this sample?

a) 0.010 c) 7.6

b) 1/9 d) 10

20. At a given temperature, gaseous ammonia molecules (NH3) have a velocity that is ____ gaseous sulfur

dioxide molecules (SO2).

a) greater than c) equal to

b) less than d) more inf. needed

21. The ratio of the average velocities of SO2(g) to CH4(g) at 300 K is

a) 1:4 c) 4:1

b) 1:2 d) 2:1

22. A sealed flask contains 1 molecule of hydrogen for every 3 molecules of helium at 20 °C. If the total

pressure is 400 kPa, the partial pressure of the hydrogen is…

a) 100 kPa c) 300 kPa

b) 200 kPa d) 400 kPa

23. A given mass of a gas occupies 5.00 L at 65 C and 480 mmHg. What is the volume of the gas at 630

mmHg and 85 C?

a) 5.00 x x

b) 5.00 x x

c) 5.00 x x

d) 5.00 x x

e) 5.00 x x

24. Which statement best explains why a confined gas exerts pressure?

a) the molecules are in random motion

b) the molecules travel in straight lines

c) the molecules attract each other

d) the molecules collide with the container walls

25. CH4 gas and O2 gas are together in a container. Which statement correctly describes the velocities of

the two molecules.

a) The two molecules have the same average velocity.b) The CH4 is moving twice as fast as the O2.

c) The CH4 is moving faster, but not twice as fast as the O2.

d) The O2 is moving faster than the CH4.

Nuclear Chemistry

3. Iodine-131 undergoes "beta decay". What other particle is produced?(A) Xe-131 (C) I-130(B) Te-131 (D) Sb-127

4. What is the charge carried by a beta particle?(A) -1 (B) 0 (C) +1 (D)+2

5. What type of radiation is simply a very energetic form of light?(A) alpha (C) gamma(B) beta (D) positron

6. Md-256 decays spontaneously with a half-life of 1.5 hours. Which one of the following statements is true about Md-256 after 3.0 hours?(A) All of the Md-256 will be decayed.(B) 75% of the Md-256 will remain.(C) 50% of the Md-256 will remain.(D) 25% of the Md-256 will remain.

7. In a decay series of Th-232, the first three steps involve an alpha decay and then two beta decays. What is the result of these decays?(A) Th-228 (C) Fr-224(B) Rn-228 (D) Pb-207

Questions 8 - 10 refer to this graph:

8. According to the above data, what is the half-life of the substance?(A) 1.0 hrs (C) 3.0 hrs(B) 2.3 hrs (D) 8.0 hrs

9. What percent of the original sample remains after 4 hours?(A) 80% (C) 60%(B) 75% (D) 40%

11. Iodine-131 has a half-life of 8 days. What percent of a sample remains after 24 days?(A) 75% (C) 50% (B) 25% (D)12.5%

12. Which of the following describes what occurs in the fission process?(A) a heavy nucleus is fragmented into lighter ones.(B) a neutron is split into a proton and an electron.(C) two light nuclei are combined into a heavier one.(D) a particle and an anti-particle turn completely into energy.

13. The measured mass of neutral Li-6 is 6.01512 amu. What is the mass defect of this isotope?mass electron 0.0005468 amumass proton 1.007277 amumass neutron 1.008665 amu

(A) 0.03435 amu (C) 0.04947 amu(B) 0.01512 amu (D) 0.03038 amu

14. One mole of H2O has a mass of 18.0 grams or 0.0180 kg. Knowing that the speed of light, c, is 3.0 x

108 m/s, calculate the energy in one mole of H2O if all of its mass were changed to energy.

[Note: 1 J = 1 kg·m2/s2](A) 1.6 x 1015 J (C) 1.8 x 1018 J(B) 5.4 x 1010 J (D) 2.0 x 1013 J

15. The "control rods" in a nuclear reactor are designed to absorb ________ and are made of the element ____(A) energy, Cd(B) uranium atoms, Pb(C) alpha particles, water(D) neutrons, Hf

16. The rate constant for decay of 218Po is 0.231 min-1. What is the half-life of this isotope?(A) 0.231 min (C) 13.6 min(B) 3.00 min (D) 4.33 min

17. Which of the following is a fission reaction?

18. Which of the following is a fusion reaction?

19. In the nuclear equation,

the letters Z and A are, respectively(A) 90 and 242 (C) 94 and 234(B) 94 and 242 (D) 90 and 234

20. Radioactive C-14 has a half-life of about 5,000 years. If a fossil is only about 6% as radioactive as expected for living tissue of the same mass, the age of the fossil is about:(A) 5,000 yrs (C) 20,000 yrs(B) 10,000 yrs (D) 40,000 yrs

21. The half-life of 210Bi is 5.0 days. What is the rate constant for decay for this isotope, with the correct units?(A) 0.20 days (C) 0.14 days(B) 0.20 days-1 (D) 0.14 days-1

22. A 10.0 gram sample of thorium-227 decays to 8.51 grams in a period of 3.00 days. What is the rate constant for this decay?(A) 0.0611 day-1 (C) 0.0851 day-1

(B) 0.0913 day-1 (D) 0.0538 day-1

23. A sample of neptunium-234, with a half-life of 4.40 days, is allowed to decay for 7.10 days. What percent of the original sample remains?(A) 19.9% (C) 30.6%(B) 61.9% (D) 32.7%

24. Cobalt-64 decays by a first order process by the emission of a beta particle. The Co-64 isotope has a half-life of 7.8 minutes. How long will it take for 15/16 of the cobalt to undergo decay?(A) 7.8 min (C) 23.4 min(B) 15.6 min (D) 31.2 min

25. Referring to the figure below, which one of these corresponds to the fission process?(A) III II(B) I II(C) III V(D) V IV

IMF’s, Liquids, and Solids1. Surface tension in a liquid is due to the fact that

a) surface molecules are pulled toward the interior

b) liquids tend toward lowest energy

c) PE is increased for molecules at the surface

d) interior molecules are attracted in all directions

e) all of the above

2. In which one of the following will dipole-dipole attractions play the most significant role as the

intermolecular attraction?a) HCl d) H2O

b) NaCl e) NH3

c) Kr

3. With which type of substances do London dispersion forces play the most significant role?

a) polar molecules d) non-polar molecules

b) metals e) network compounds

c) ionic compounds

4. The heat of vaporization of H2S, at its boiling point (–61°C) is 18.8 kJ/mol. What mass of H2S can be

vaporized (at its boiling point) with 100 kJ of energy?

a) 100 x d) 18.8 x

b) 34.1 x e) 100 x

c) 61 x 18.8 x 100 x 34.1

5. Which one of the following substances exhibits the strongest intermolecular forces of attraction?a) CH4 d) CH3OH

b) C2H6 e) CH3Cl

c) C3H8

6. For which substance would you predict the highest heat of vaporization?a) F2 b) H2O c) HF d) NaCl e) Br2

7. Which of the following will change the equilibrium vapor pressure of a liquid?

I. Heat up or cool down the liquid

II. Increase the Volume of the container

III. Change the pressure above the liquid

a) I only d) I and III only

b) I and II only e) II and III only

c) I, II, and III

8. Which of the following statements describes a substance above its critical point?

a) the substance can be liquefied

b) the vapor and liquid phase become indistinguishable

c) the substance experiences no intermolecular interactions

d) there is a distinct phase boundary between the liquid and vapor

e) all of the above

9. At what temperature will the liquid (whose vapor pressure is shown below) boil if the air pressure is

reduced to 380 mmHg?

a) 30°C b) 50°C c) 70°C d) 100°C

e) the liquid will not boil at this pressure

10. Which one of the following is linked with the correct intermolecular force of attraction?a) NH3....................dipole-dipole

b) AlH3...................London dispersion forces

c) H2.......................hydrogen bonding

d) C2H4...................covalent bonding

e) HCl.....................ionic

11. The vapor pressure graph of an unknown liquid is shown below. Which of the following statements

about this liquid is/are true?

I. This liquid has weaker IMF’s than water.

II. The liquid’s normal boiling point is around 75°C.

III. The liquid boils at room temperature when the pressure is dropped to about 0.25 atm.

a) II and III only d) I only

b) II only e) I, II, and III

c) I and III

12. How much energy does it require to melt 25.0 g benzene, C6H6? The heat of fusion of benzene is 2.37

kJ/mol. [molar mass = 78.0 g/mol)

a) 8.25 kJ d) 0.759 kJ

b) 59.3 kJ e) none of these

c) 4625 kJ

13. What type of solid(s) can contain covalent bonds?

a) molecular d) network

b) metallic e) all but “b”

c) ionic

14. Which type of solid generally has the highest melting point?

a) metallic c) molecular

b) ionic d) network

15. Which substance below exhibits the weakest IMFs?a) IF3 b) SO2 c) CO2 d) SiO2 e) PH3

16. During the condensing of a liquid, the kinetic energy _____ and the potential energy ______.

a) stays the same, increases

b) increases, decreases

c) increases, increases

d) decreases, stays the same

e) stays the same, decreases

17. The phase diagram of a substance is given below. What occurs when the substance is heated from

100° C to 120 °C at 3 atm pressure?

a) it melts d) it freezes

b) it sublimes e) no phase

c) it boils change occurs

18. A typical phase diagram for a substance is given below. At what point on the diagram do solid and

liquid exist at equilibrium?

a) A b) B c) C d) D e) E

19. Which one of the following as solids has a crystal structure containing discrete (separate) molecules?

a) potassium d) carborundum, SiC

b) glass e) hydrogen

c) quartz

20. The heat of sublimation of a compound equals

a) heat of fusion plus heat of vaporization

b) heat of ionization plus heat of crystallization

c) heat of vaporization minus heat of fusion

d) heat of vaporization plus heat of crystallization

e) heat of crystallization plus heat of vaporization

21. The normal boiling point of a liquid

a) is 100 °C at 1 atm pressure.

b) is the temperature at which the vapor pressure is 1 atm.

c) is the temperature at which liquid and vapor are in equilibrium.

d) is the temperature at which the vapor pressure equals the external pressure.

e) is the temperature at which there is a continuous formation of gaseous bubbles in the liquid.

22. The vapor pressure of a liquid increases with an increase of temperature. Which of the following best

explains this increase?

a) The average kinetic energy of molecules is greater, thus more molecules can enter the gaseous

state.

b) The number of gaseous molecules above the liquid remains constant but these molecules have

greater average kinetic energy.

c) the faster moving molecules in the liquid exert a greater pressure.

d) All the molecules have greater kinetic energies.

e) The intermolecular forces between the molecules becomes less at higher temperatures.

23. Which of the following indicates very strong intermolecular forces of attraction in a liquid?

a) A very low boiling point.

b) A very low critical temperature.

c) A very low heat of vaporization.

d) A very low vapor pressure.

e) A very low surface tension.

24. The compounds Br2 and ICl have almost identical molecular weights, yet ICl boils at 97°C and Br2

boils at 59 °C. The best explanation for the difference isa) ICl is an ionic compound and Br2 is covalent.

b) ICl is a nonpolar molecule and Br2 is polar.

c) ICl has a longer bond than that in Br2 .

d) ICl has a measurable dipole moment (is polar) and Br2 does not (is nonpolar).

e) ICl has a stronger bond than that in Br2 .

25. In some compounds the hydrogen atom is covalently bonded to one atom and simultaneously attracted

to another atom in another molecule by an electrostatic interaction. This interaction can occur when

hydrogen is bonded toa) Cl b) Si c) N d) C e) Br

26. Which of the following compounds shows an abnormal boiling point due to hydrogen bonding?

a) CH3NH2 d) CH3Cl

b) CH3OCH3 e) HCl

c) CH3SH

27. Which of the following has the lowest boiling point?

a) H2O d) H2Te

b) H2S e) NH3

c) H2Se

28. Which of the following would be expected to have the highest heat of vaporization?

a) H2O c) HF

b) NH3 d) all three are the same

29. Which element is considered a covalent/network solid?a) Cr b) O c) Xe d) B e) Na

30. Which one of the following compounds has intermolecular forces different than the others?a) quartz, SiO2 d) C(graphite)

b) C(diamond) e) silicon carbide, SiC

c) carbon dioxide, CO2